H. A. LAITINEN AND C. J. NYMAN have presented evidence of ionic bonding for sulfate and secondary phosphate ions in the second sphere. Results of their work were not conclusive with other electrolytes. It might be noted that a possible explanation based upon displacement of part or all of the coordinating groups of the original complex by electrolyte ions was ruled out for a number of reasons. First, ammonia is known to form a very stable covalent bond to cobalt (111), whereas sulfato complexes are in general much less stable. Second, from the point of view of super-complexing, which is electrostatic in character, the divalent ions such as the sulfato ion would be expected to show the clustering behavior more readily than univalent ions such as chloride, nitrate, or perchlorate, a prediction which is in accord with the experimental result. Third, the color stability of the solutions provides a reason for discounting a theory based upon displacement of the ammonia molecules in the original complex, inasmuch as displacement of coordinating groups might reasonably be expected to cause a change of color; actually no such color change occurs. Fourth, the size and charge of the ion that would result from such a displacement should lead to less hydrated ions; for example, if a sulfate ion replaced two ammonia molecules, the resulting ion would be a univalent cation which would be less highly hydrated [CONTRIBUTION FROM THE ?;OYES
CHEMICAL
Vol. 70
than the original trivalent ion and, hence, would be expected to diffuse more rapidly; if two sulfate ions replaced four ammonia molecules, a univalent negative ion would result, again with a higher diffusion rate, which would orient differently with respect to the dropping mercury electrode and might be expected to be more easily reduced if indeed any change would result; the experimental result is in the opposite direction.
summary 1. Capillary active substances such as gelatin and octyl alcohol were found to shift the half-wave potentials for the first reduction step of the hexamminecobaltic ion to more negative potentials. 2. The optimum drop rate for most favorable conformity with the Ilkovic equation was established. 3. Supporting electrolytes of good coordinating ability were found to shift the half-wave potential for the first reduction step of the hexamminecobaltic ion to more negative potentials and to decrease the diffusion rate of the cobaltic complex, A “super-complex” of anions of the indifferent electrolyte clustered about the central complex by electrostatic and ion-dipole attraction is postulated. URBANA, ILLINOIS RECEIVED APRIL12, 1948 LABORATORY OF THE
UNIVERSITY OF ILLINOIS]
Polarography in Liquid Ammonia. 11. The Electron Electrode BY H. A. LAITINEN AND C. J. NYMAN’~
The electrode reactions which might occur a t a cathode on electrolysis of anhydrous liquid ammonia solutions have been summarized by Makishima.2 For a platinum electrode, metal deposition, reduction in valence state of the metal ion, hydrogen discharge, and electron dissolution are possible processes. The dissolution o$ electrons was observed by several investigators when a salt solution, whose ions were non-reducible, was electrolyzed with a platinum cathode. On electrolysis of solutions of tetramethylammonium chloride and hydroxide and tetraethylammonium chloride between platinum electrodes at -34O, Palmaer3 obtained in each case a blue solution, which he attributed to the electron or to the “free alkylammonium radical.” Similar observations were made by Kraus4 and by Schlubach.s Forbes and Nor(1) Presented hefore the Division of Physical and Inorganic Chemistry a t the Chicago Meeting of the American Chemical Society, 70, 2241 (1948). April, 1948. For paper I, see THIS JOURNAL, (la) Present address: Department of Chemistry, The State College of Washington, Pullman, Wash. (2) S. Makishima, J. Faculty Eng. Tokyo Imp. Univ., 21, 115 (1938). (3) W. Palmaer, 2.Elektrochcm.. E, 729 (1902). (4) C. A. Kraus, THISJOURNAL, 86, 1732 (1913). (5) H.H,Schliibach, Bcr., 58, 1689 (1920).
ton6 have prepared solutions of tetramethyl-, tetraethyl-, tetrapropyl-, tetrabutyl-, trimethylbutyl-, triethylbutyl-, tripropylbutyl-, tributylmethyl-, tributylethyl- and diethyldibutylammonium ions and electrons by electrolysis of the iodides of these ions between platinum electrodes. Cady’ and Kraus* found on electrolysis of sodium solutions between platinum electrodes that the transfer of electrons from the electrode to the solution was the only cathodic process which occurred and that the opposite transfer was the only anodic process which occurred. All the processes listed for a platinum cathode can also occur a t a dropping mercury cathode. In addition, there is the possibility of amalgamation of the cation of the salt solution, as is known to be the case with the alkali metal ions.’ For electrons to enter solution from the surface of a dropping mercury cathode, the potential of electron dissolution must be more positive than the amalgamation potential of the cation of the salt solution. (6) G . S. Forbes and C. E. Norton, THIS JOURNAL,LE, 2278 (1926). (7) H . P . Cady, J . Phys. Chem., 1,707 (1897). (8) C. A. Kraus, THISJOURNAL, SO, 1323 (1908); 88,864 (1914).
Sept., 1945
ELECTRON ELECTRODE I N LIQUID AMMONIA POLAROGRAPHY
When a transfer of electrons from the electrode to the solution, or the reverse process, takes place, the electrode can be considered to be an electron electrode. It was found possible to study these processes polarographically with both platinum and dropping mercury electrodes. Experimental The electrolysis cells and apparatus for preparing the anhydrous liquid ammonia solutions, as well as the thermostat, have been described previously.1 Unless otherwise specified, the cell with the internal reference electrode was used. The dropping mercury electrode had the following characteristics in a saturated solution (0.0057 M ) of tetrabutylammonium iodide in liquid ammonia a t -36". At a pressure of 20 cm. of mercury, the drop time t was 5.1 sec. (open circuit), and the mass of mercury flowing through the capillary was 1.184 mg./sec. The platinum electrode was of a type described by Laitinen and Kolthoff,; and consisted of a platinum disc sealed t o the end of a glass tube in such a manner that the disc was vertical when the electrode was in place. The area of the electrode was about 3.1 sq. mm. The reference electrode used in this investigation was a mercury pool in 0.005 M iodide solution. A Sargent Model XX polarograph was used. The calibration of this instrument was checked by inserting a known resistance in place of the electrolysis cell and measuring accurately the applied potential. With these data the current flowing through the circuit was calculated and compared with that indicated by the recorder. All applied potential values were checked by means of a student type potentiometer. The ammonia, obtained from the Matheson Company, was dried by condensation on sodium before being distilled into the electrolysis cell. The alkali metal salts were C. P. materials of commerce and were dried a t 110' before use. The tetraalkylammonium iodides were prepared in an organic preparatibns course a t the University of Illinois. The tetramethyl- and tetraethylammonium iodides were recrystallized from 95yc ethanol, the tetrapropylammonium iodide from 50oJo ethanol-50% ethyl acetate solution, and the tetrabutylammonium iodide from absolute ethyl acetate. Residual current curves of 0.005 M solutions of thest: salts in liquid ammonia showed no waves a t potentials which corresponded to alkali metal reduction, and the salts were therefore considered free of such impurities. Sodium amide was prepared by the reaction of sodium with liquid ammonia in the presence of a trace of ferric nitrate hexahydrate as described by Greenlee and Henne.Io
Theoretical A polarogram made with a dropping mercury electrode on a 0.001 molar solution of sodium iodide in a saturated solution of tetrabutylammonium iodide is shown in Fig. 1. It is proposed that the increase in current a t potentials more negative than 2.3 volts was due t o the dissolution of electrons. Assuming, for the present, that this was the process which occurred, an equation for the current-voltage curve obtained can be derived. The potential of an electrode behaving as an electron electrode should be given by a Nernst expression of the type (9) H. A. Laitinen and I. M. Kolthoff, THIS JOURNAL, 81, 3344 (1939). (10) K. W. Greenlee and A. L. Heqne, "Inorganic Syntheses," 2, 128 (1946).
3003
where Ee is the potential of the electrode; E: is the standard potential of the electron electrode; and C, is the concentration of the electron a t the surface of the electrode in moles per liter. Assuming that the Ilkovic equation describes the behavior of an electron with regard to its diffusion from a dropping mercury electrode, then C. = i/6.05 X 105D1/~mp/lt%
(2)
where i is the current which flows (microamperes); D is the diffusion coefficient of the electron; rn is the mass of mercury flowing through the capillary (mg./sec.) ; and t is the drop time of the capillary in seconds. On substitution of C, from equation 2 into equation 1, the expression E. = Eo
is obtained. writing
- gF l
n il6.05 X 1 0 ~ D 1 / ~ m p / ~ t(3) l/~
Equation 3 can be simplified by
a=A -
In i
(4)
In 6.05 X 1O6D1/:ma/at'/6
(5)
where A =
+
Equations 3 and 4 give the potential of a dropping mercury electrode when a current of i microamperes flows due to the dissolution of electrons. The constant A represents the potential of the electrode when a current of one microampere is flowing. A depends on E", and on the characteristics of the electrode in a manner given by equation 5. If the process is reversible, a plot of E versus log i should have a slope of -0.047 a t -36'. The value of A can be determined most accurately from a plot of this type.
A
-1.4
-1.6
-1.8
-2.0
-2.2
-2.4
E. Fig. 1.-Polarogram of 0.001 M NaI in saturated (C4Hs)dNI, dropping mercury electrode.
An expression similar to equation 4 would also describe the current-voltage curve obtained when using a platinum electrode. The constant A would change for different electrodes, because the diffusion of electrons would vary with the area and
H. A. LAITINEN AND C. J. NYMAN
3004
shape of the electrode and with the orientation of the electrode surface.
Vol. 70
4.4 X lo-'' mole per liter in a neutral solution. If the amide ion concentration were increased to
mole per liter, the ammonium ion concentration would be decreased to 1.9 X mole per On electrolysis of solutions of uni-univalent liter. From the Nernst equation, this would salts with a dropping mercury electrode, the three cause hydrogen discharge to be more difficult and possible cathodic reactions which must be con- the constant A to be 0.58 volt more negative than sidered are hydrogen discharge, amalgamation of in a neutral solution. A polarogram was run on a mole of sodium the cation, and electron dissolution. Since it has solution which contained been previously reported that the alkali metal ions amide per liter and which was saturated with can be reduced polarographically in liquid am- tetrabutylammonium iodide. In this solution the monia, salts of these ions cannot be used as elec- value of A was found to be 0.70 volt more negative trolytes in a study of mercury as an electron elec- than the sodium half-wave potential. Since there trode. Tetrabutylammonium iodide was chosen was no shift in the value of A with a decrease in as the one to be studied in this connection because ammonium ion concentration, the process obit gives a very negative break in water solution, served could not have been hydrogen discharge. If the increase in current was due to reduction of and it was hoped that i t would prove to be nonreducible in liquid ammonia even a t the most the cation to form an amalgam. a limiting current negative potentials. From the polarogram in should have been observed, such as the one obFig. 1i t is seen that the value of A is 0.72 volt more tained on the polarographic reduction of 0.005 monegative than the sodium half-wave potential. lar sodium iodide solution (Curve IV, Fig. 2). In the discussion which follows, it will be shown On the other hand, if no amalgamation occurred, that this break is caused by the transition of elec- the only other cathodic process which could have trons from the electrode to the solution and not by occurred was the dissolution of electrons. A polarogram for a saturated solution (0.0057 -44) of one of the other two processes. If the increase in current a t potentials more tetrabutylammonium iodide (Curve 11, Fig. 2) negative than 2.3 volts was caused by hydrogen showed no limiting current even at potentials as discharge, then the value of A should be a function negative as 4 volts. This was a good indication of ammonium ion concentration. Pleskov and that the process was not amalgamation of the Monossonll found the dissociation constant of tetrabutylammonium ion. If the process was that of electron dissolution, liquid ammonia to be 1.9 X a t -50'. Assuming it to be the same a t -36') the ammonium then the cation of the salt solution should have no ion and amide ion concentrations are equal to effect on the potential of electron dissolution, and, therefore, the tetrabutylammonium ion may be replaced by any ion which is not reduced a t these potentials. In this connection, a series of experiments was carried out using tetramethyl-, tetraethyl-, tetrapropyl- and tetrabutylammonium iodides. A polarogram (Curve V, Fig. 2 ) of a saturated solution (0.0042 M ) of tetramethylammonium iodide showed a limiting current similar to that shown by sodium ion. This was not unexpected since I M c C O ~and ~ ~co-workers .~~ had previously prepared tetramethylammonium amalgam a t - 34' by electrolysis of the chloride salt in absolute methanol. They were unable to prepare tetraethylammonium amalgam under similar conditions. A polarogram (Curve I, Fig. 2) of a solution of tetraethylammonium ion showed no limiting current, but the constant A was more positive than for the tetrabutylammonium iodide solution. This can be seen more conveniently from Fig. 3, which is plotted on a larger scale. The tetraethylammonium iodide curve was very erratic and not nearly so well defined as the tetrabutylammo-2.0 -3.0 -4.0 nium iodide curve. I t appears possible that both E. electron dissolution and amalgamation of the caFig. 2.-Polarograms with dropping mercury electrode: tion might have occurred with the tetraethylamcurve I, 5 X M (CzH5)4NI; curve 11, satd. (C4Hs)rNI; monium iodide solution. The current-voltage curve 111, 5 X M (CaH∋ curve IV, 5 X 10-8 curve (Curve 111, Fig. 2 ) for the tetrapropylM NaI; curve IT, satd. (CH&NI.
Data and Discussion
(11) V. A. Pleskov and A. hl. Monosson, Acta Physicochim.
S. 8.. 1, 713 (1935).
U.R.
(12) H. N. McCoy and W. C. Moore, THISJOURNAL, 33, 273 (1911). (13) H. N. McCoyand F. L. West,J. Phys. Chem., 16,261 (1912).
ELECTRON ELECTRODE IN LIQUIDAMMONIA POLAROGRAPHY
Sept., 1048
-2.27
-2.31
-2.35
-2.39
3005
-2.43
E. Fig. 3.-Polarograms with dropping mercury electrode: curve I, 5 X 114 (C:Hj)rNI; curve 11, 5 X M (CaH7)4KI; curve 111, satd. ( C ~ H Q ) ~ N I .
ammonium iodide was almost identical in every respect with the current-voltage curve of tetrabutylanimonium iodide. This curve is also shown in Fig. 3. The values of A obtained from these curves by plots of E versus log i are recorded in Table I along with the slopes of such plots. The fact that these quantities are almost identical for tetrapropyl- and tetrabutylammonium iodide solutions and the fact that no limiting current was obtained indicated that the same process, electron dissolution, took place in both solutions. It might be expected that the current-voltage curves for sodium ion and tetramethylammonium ion would show a sharp increase in current a t -2.36 volts E i that is the potential a t which electron dissolution occurs. One reason that a sharp rise is not encountered is the fact that the concentration of positive ions a t the surface of the electrode is vanishingly small, because they are being amalgamated as fast as they reach the electrode surface. If the cations are being amalgamated, then the solution surrounding the electrode would have a very high resistance as compared to the bulk of the solution, and the flow of current would be impeded. From the slopes of the currentvoltage curves a t the more negative potentials] the maximum resistances of the solutions a t the electrode surface can be calculated. The maximum resistances of the solutions of 0.005 tetraethyl-, 0.005 d l tetrapropyl- and saturated tetrabutylammonium iodide solutions were about 5,000 ohms, and the resistances of saturated tetramethyl ammonium iodide and 0.005!J sodium iodide solutions were about 20,000 ohms. A gradual rise in current is observed a t very negative potentials for the latter two salts on electrolysis, and this must correspond to some cathodic process which in all probability is electron dissolution. It has already been mentioned that liquid ammonia solutions of cations and electrons can be prepared electrolytically using a platinum cathode, if the cations are stable in the presence of the solvated electron. As in the case of the dropping mercury electrode, changing the type of cation of
-2.28
-2.32
-2.36
-2.40
E. Fig. 4.-Polarograms
with stationary electrodes: curve
I, satd. (C4Hs)dNI with mercury surface; curve 11, 5 X Pt,I , lo+ M NaI, Pt surface; curve 111, satd. ( C ~ H Q ) ~ N surface; curve IV, 5 X 10-3 M LiI, Pt surface; curve V, M Na, Pt satd. (CH&NI, Pt surface; curve VI, 6 X surface.
the salt solution should have no effect on the potential of electron dissolution. Polarograms of solutions of lithium, sodium, tetramethylammonium and tetrabutylammonium iodides using a platinum electrode are shown in Fig. 4, and the similarity of the curves bears out the fact that the process is the same in each case. The values of A were nearly constant] but the slopes of the E versus log i plots differed from the theoretical value by a considerable margin. These quantities are also recorded in Table I. The difference between the observed slopes and the theoretical ones may be due to poor diffusion conditions in the solutions near the electrode surface or possibly to a potential drop caused by the resistance of the solution. In any event, diffusion to a stationary electrode is not so well defined as to a dropping mercury electrode. TABLE I Electrode
A , v.
Drop Hg Drop Hg Drop Hg Pt Pt Pt Pt Stationary Hg Pt
-2.26 -2.34 -2.36 -2.35 -2.34 -2.35 -2.34 -2.34 -2.35
Slope E log i
91s.
..... -0.064 - .063
-
-
.lo4 .120 ,105 ,098 .093 .073
To determine whether the electrode material affected the potential of electron dissolution, a polarogram was run on a tetrabutylammoniurn
3006
H. A. LAITINEN AND C. J. NYMAN
Vol. 70
iodide solution using a stationary mercury elec- be a corresponding anodic process with the transtrode. The platinum electrode used in the experi- fer of electrons from the solution to the electrode. ments described previously was plated with cop- The anodic transfer can take place a t any potenper and then with mercury until a smooth film of tial more positive than that of electron dissolution. mercury covered the electrode surface. If the On electrolysis of a solution of sodium metal, an same process occurred a t the mercury surface as anodic current was observed a t the platinum elecat the platinum surface on.electrolysis of a tetra- trode, and the polarogram obtained is shown by butylammonium iodide solution, then the values Curve VI in Fig. 4. An H-type cell, with a sinof A and the slopes of the E versus log i plots tered glass disc to separate the anode and cathode should be comparable, because the shape and area compartments, was used for this experiment. of the electrodes were approximately the same. One compartment contained the sodium solution This was observed experimentally and, therefore, and the platinum electrode, and the other conthe process was the same. These experiments tained the mercury reference electrode in a satualso indicate that there is no detectable overvolt- rated solution of tetrabutylammonium iodide. A age to the dissolution of electrons either from a trace of mercuric iodide was added to the tetraplatinum or mercury surface. Curve I in Fig. 4 is butylammonium iodide solution to depolarize the the polarogram of a tetrabutylammonium iodide reference electrode. The potential a t which the solution using the stationary mercury electrode, anodic current has decreased by one microampere and the data obtained from it are recorded in in this experiment should be the same as the potential a t which the cathodic current has increased Table I. The platinum electrode used above was plated by one microampere on electrolysis of a salt soluwith copper and, as in the case of the stationary tion. The value of A observed for the sodium mercury electrode, it should have the same area solution was very close to that observed for salt and shape as the platinum electrode. It was solutions as can be seen from Table I. The fact found on electrolysis of a saturated solution of that the curve goes smoothly through the zero curtetrabutylammonium iodide using the copper elec- rent line indicates that there is no overvoltage for trode that A was more positive than for a platinum either the dissolution process or the reverse procelectrode, and that it varied between -1.5 and ess. -1.8 volts. Reproducible values could not be An unsuccessful attempt was made to obtain an obtained. Apparently, hydrogen discharge oc- anodic curve using the dropping mercury electrode curred a t the copper surface more readily than in a sodium solution. The failure to obtain a the dissolution of electrons. This would indicate polarogram was attributed to the fact that sodium that freshly plated copper has a low hydrogen solutions are not stable in the presence of mercury overvoltage in liquid ammonia, while bright plati- because of sodium amalgam formation. num and mercury have high hydrogen overvoltA solution of tetrabutylammonium ions and ages. Pleskov14 has reported that mercury has a electrons was prepared by electrolyzing solutions higher hydrogen overvoltage in liquid ammonia of tetrabutylammonium iodide in an H-type cell between mercury pool electrodes. Anodic curthan in water. Platinum black electrodes have been reported rent-voltage curves were obtained for solutions by Pleskov and Monossonl6 to function as rever- prepared in this manner, but the results were not sible hydrogen electrodes in liquid ammonia. reproducible. An anodic current was observed a t This indicates that such electrodes have little or no potentials as negative as -2.3 volts. hydrogen overvoltage. When a platinum black Makishima2 calculated the standard potential electrode was used as the cathode for obtaining a of the electron electrode to be 2.45 volts more current-voltage curve of a solution of tetrabutyl- negative than the standard potential of a lead elecammonium iodide, a large current was observed as trode or 2.04 volts more negative than the standsoon as a negative potential was applied. At a ard potential of the hydrogen electrode a t -34'. constant potential, 0.1 volt more negative than Forbes and Nortons have measured the oxidation the reference electrode, the current flowing slowly potentials of the solutions of tetraalkylammonium decreased. This would be the state of affairs if ions and electrons mentioned earlier as well as the the process were hydrogen discharge from an un- oxidation potentials of solutions of lithium, sobuf€ered solution. That hydrogen discharge is dium and potassium. The concentrations of catnot the process which takes place a t the surface of ions and electrons in these solutions were of the a bright platinum electrode is indicated by the fact order of 0.005 M , and the oxidation potentials of that solutions of electrons can be prepared by all were found to be 2.590 * 0.015 volts more electrolysis using platinum cathodes and that negative than the potential of a silver electrode in metal solutions are stable in the presence of bright a saturated solution of silver nitrate a t -778'. platinum electrodes. There was no general trend in the variation of poI n addition to having a cathodic process with tential with concentration of electron, and it may the formation of an electron solution, there should be possible that their analytical method for deter(14) V. A. Pleskov, Acfo Physicochim., U.R.S.S.,11,305 (1939). mining the concentration of electron was not accu(15) V. A. Pleskov and A. M. Monosson, ibid, U.R. S. S.. 1, 871 rate enough to establish such a trend. (1935).
Sept., 1948
ELECTRON ELECTRODE IN LIQUIDAMMONIA POLAROGRAPHY
By a consideration of the current-voltage curve for a solution of sodium metal, the standard potential of the electron electrode can be evaluated. The potential at which the current-voltage curve crossed the zero current axis is a function of the surface concentration of electron, and since a t this potential no current is flowing, the concentration a t the surface of the electrode should be equal to the concentration in the bulk of the solution. The current-voltage curve crossed the zero current axis at -2.388 volts, and the concentration of electron at the electrode surface was 0.006 mole per liter. On substitution of these data into equation 1, the standard potential of the electron electrode was estimated to be 2.49 volts more negative than the mercury reference electrode. It is possible to relate this potential to the standard potential of the hydrogen electrode in liquid ammonia. The difference between the potential of the mercury reference electrode and the potential of a lead electrode in an 0.1 N lead nitrate solution was previously reported' as 0.318 volt a t -36'. Pleskov and MonO S S O have ~ ~ ~ reported the difference of potential between a lead electrode in 0.1 N lead nitrate solution and a hydrogen electrode in 0.1 N ammonium nitrate to be 0.352 volt a t -50'. The partial pressure of hydrogen in the system was 450 mm. Neglecting the temperature coefficient of the above cell and the activity coefficient of the ammonium ion, the difference of potential between a lead electrode in 0.1 N lead nitrate solution and the standard potential of the hydrogen electrode was calculated to 0.310 volt. Combining these cells, the standard potential of the electron electrode is 1.86 volts more negative than the standard potential of the hydrogen electrode. Equation 5 can be used to calculate the standard potential of the electron electrode, since the remaining terms can be evaluated. The diffusion coefficient can be calculated from the equivalent conductance of the electron at infinite dilution, and the other terms are experimental quantities. Kraus found that the limiting conductance of the electron a t infinite dilution was 910 ohm-1-cm.2equiv.-'. Using the equation D = RTXO/zP (6) derived by N e r n ~ t , ' ~ the . ~ ' diffusion coefficient D of the electron can be calculated. R is 8.317 volt-coulombs, T is the absolute temperature, Xo is the equivalent conductance of the electron a t infinite dilution, z is the charge of the electron, and F is 96,500 coulombs. At -36" D is 1.93 X sq. cm./sec. The drop time of the electrode was 1 sec., and m was 1.184 mg./sec. For polarograms of saturated solutions of tetrabutylammonium iodide, the average value of the constant A was -2.36 volts with reference to the potential of a mercury pool in a saturated solution of tetrabutylammonium iodide. Substituting these data (16) W.Nernst, 2:. physik. Chem., 2, 613 (1888). (17) J. J. Lingane and I. M. Kolthoff, THISJOURNAL, 61, 826 (1939).
3007
into equation 5, the standard potential of the electron electrode was estimated to be 2.55 volts more negative than the mercury reference electrode and 1.92 volts more negative than the standard potential of the hydrogen electrode a t -36'. These two values for the standard potential of the electron electrode are in fair agreement with the theoretical values calculated by Makishima. The potential a t which the dissolution of electrons occurs has an important theoretical and practical significance. It represents the limit to the negative potentials which can be applied to an electrode for polarographic measurements in liquid ammonia. It is possible to relate the standard potential of the electron electrode to the solubility of metals in the following way. When a metal dissolves in liquid ammonia to form solutions of cations and electrons, according to the reaction M
--+
+ ne-
MnC
an anodic and cathodic process are taking place simultaneously. In the dissolution of metal ions, the potential of the metal is given by the expression RT El = PM - In C p + (7) nF For the dissolution of electrons, the potential of the metal is given by the expression
+
RT E z = E O , - - nF In (Ce-)"
(8)
Since both processes are taking place a t the same electrode, the potential E1 must be equal to Ez at all times, and therefore
pM - E",
=
- gn F
~ Cp+(Ce-)n n
(9)
The logarithmic term Cw+(C,-)" is equivalent to the solubility product of the metal, and the solubility of the metal is given by CM~+. These relationships have been pointed out previously by Makishima2 and were used by him in evaluating the standard potential of the electron electrode. Equation 9 implies that univalent metals whose standard electrode potentials are more negative than that of the electron electrode will be more soluble than one mole per liter. Conversely, univalent metals whose standard potentials are more positive than the standard potential of the electron electrode have a maximum solubility less than one mole per liter. Equation 9 has been found to be valid, a t least qualitatively, for those alkali metals and alkaline earth metals whose solubilities and standard electrode potentials are known. Theoretically all metals would be soluble to an extent governed by equation 9. The most positive metals would have an infinitesimally low solubility. Summary 1. Platinum and mercury electrodes were found to function as electron electrodes in liquid ammonia solutions when the cation of the in-
3008
H. G. THODE, J. MACNAMARA, F. P. LOSING AND C. B. COLLINS
different electrolyte was non-reducible. An equation was derived for the polarograms obtained and for estimation of the standard potential of the electron electrode from polarograms of salt solutions. 2. Two different types of measurements gave values of - 1.86 and - 1.92 volts for the standard potential of the electron electrode versus the stand-
[CONTRIBUTION FROM THE
Vol. 70
ard potential of the hydrogen electrode at -36’. These values are roughly in agreement with a theoretical value of -2.04 volts calculated by Makishima. 3. A relationship between the solubility of metals in liquid ammonia and the standard potential of the electron electrode was presented. URBANA, ILL.
RECEIVED APRIL15, 1948
DEPARTMENT OF CHEMISTRY, MCMASTER UNIVERSITY]
Natural Variations in the Isotopic Content of Boron and its Chemical Atomic Weight BY H. G. TIIODE,J. hfACNAMARA, F. P. LOSINGA N D
c. B. C O L L I N S
Introduction ally a t normal temperatures. Of particular interSeveral years ago it was suggested by W. J. est is the predicted difference in vapor pressures of Whitehouse and B. Pontecorvo that small varia- isotopic boron chlorides. These differences pretions in the isotopic abundances of boron might ex- dicted in the thermodynamic properties of isotopic plain discrepancies in neutron capture crossection boron compounds suggest that some isotopic exmeasurements made by different investigators change process might OCCUT in nature which would using boron as a standard. As a result of this sug- result in a slight fractionation of the isotopes. A change in the ratio of the boron isotopes can, gestion, seven or eight boron mineral samples were collected from various parts of the world with the of course, be effected by neutron irradiation. Howobject of determining the variation, if any, in the ever, neutron intensities in the earth’s surface due abundances of the boron isotopes with a mass to cosmic ray and other sources would seem a t the spectrometer. However, interest in the neutron present time to be too low, especially in geological capture crossection aspect of the problem de- deposits below the surface, to account for possible creased and the problem was dropped for a while variations in the isotopic ratios. To date, nine because of other work. The problem was again samples of boric acid, eight of known origin,.pretaken up during the past year because of general pared from boron minerals obtained from widely interest in the abundances of the boron isotopes. separated geological deposits, have been investiInghraml made a careful study of the isotopic con- gated. These samples were converted to boron tent of boron, but unfortunately did not publish in- trifluoride gas and the ratio B1l/Bl0 determined formation on the sources of his samples. Other with a mass spectrometer. The results obtained investigators have demonstrated variations in the for the nine samples are reported in this paper. isotopic content of carbon,2and oxygen3in nature, Experimental and it would not be surprising to find some variPreparation of Samples.-The boron trifluoride samples ations in the isotopic content of boron. used in the mass spectrometer investigations were preUrey and Cireiff first showed from thermodyna- pared by a method described previously.6 Stoichiometric quantities of potassium fluoborate, anhydrous boric oxide mic considerations that the isotopes of the light and concentrated sulfuric acid are mixed together in a elements did differ slightly in chemical properties. flask and the mixture is heated strongly in vucuo. The These differences in chemical properties result in boron trifluoride produced is bubbled through concensome fractionation of the isotopes in natural proc- trated sulfuric acid saturated with boric oxide to remove hydrogen fluoride and moisture. The boron trifluoride esses. Isotopic exchange constants involving the is condensed by means of liquid air and is further purified isotopes of carbon, nitrogen, oxygen and sulfur by distilling from one ampoule t o another, discarding the have been calculated by the methods of statistical first and last portions. Care was taken a t this point to avoid possible fractionation of the isotopes. mechanics and checked by experiments. The anhydrous boric oxide used is prepared by dehyRecently IJrey6 made a further study of the drating boric acid in a vacuum oven for two hours a t thermodynamic properties of isotopic substances 200” and the potassium fluoborate is prepared by treating and reviewed the theoretical and experimental re- boric acid with hydrogen fluoride to produce fluoboric sults obtained to date. The boron exchanges acid which is converted to form the potassium salt by titration with 5 N potassium hydroxide. The boron which have been investigated from a theoretical samples came in the form of boric acid. All of these but point of view are limited to the boron halides. one were prepared from borates by the usual treatment Exchange processes involving these particular with sulfuric acid. The potassium salts are crystallized compounds cannot readily be realized experiment- out first a t elevated temperatures, after which the tem(1) Inghram, Phys. Reo., 70, 6.53 (1946). (2) Nier and Gulbransen, THISJ O U R N A L , 61, 597 (1939). (3) Murphy and Nier, P h y s . R P Z ’69, . , 771 (1941). (4) Urey and Greiff, THISJ O U R N A L , 67, 321 (1935). ( 5 ) Urey, J. Chr!m. Suc., Part 2, 662 (1947).
perature is reduced to bring out the borax. (The borax is then treated with sulfuric acid t o produce boric acid.) The boron sample obtained from Tuscany in Italy is ob( 6 ) “Inorganic Syntheses,” Vol. I, McCraw-Hill Book Co.,New York, N. Y.,1939,p. 21.
